Inorganic Chemistry
By
Dr. Khalil K. Abid
Lecture 12
The Covalent Bonding
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The Covalent Bond
Generally it forms between the non metallic elements of the periodic table. It forms by electron sharing
There are several other nonmetal
elements that also form diatomic
molecules:
• F2, Cl2, Br2, I2, O2 and N2 and others
that form polyatomic molecules
• P4 , S8 and Se8
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The Covalent Bond Model
Covalent bonding leads to a greater electron density between the atoms
• In a covalent bond, electrons from two atoms are shared.
• There are rules to how covalent bonds form.
Like ionic bonding, covalent bonds form as a way of giving each atom 8 valence electrons (except
hydrogen, which only wants 2 valence electrons).
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Properties of the compounds with covalent bonds: Formed by a system of continuous covalent
bonds. Non conductive LATTICES both in the solid and in the molten state.
• Generally have low melting points and boiling points compared to ionic compounds.
Diamond, boron nitride, quartz (SiO2), silicon carbide (SiC)
1. Hard and incompressible
2. Tf high, non volatile
3. insoluble
Graphite
Diamond
Substances with covalent bonds
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Covalent Bonding in H2
:
H.
.H
H:H
Two hydrogen atoms, each with 1 electron, can share those electrons in a covalent bond.
• Sharing the electron pair gives each hydrogen an electron configuration analogous to helium.
Covalent Bonding in F2
Two fluorine atoms, each with 7 valence electrons, can share those electrons in a covalent bond.
• Sharing the electron pair gives each fluorine an electron configuration analogous to neon.
To gain a deeper understanding of the covalent bond, or shared electron bond, let's look at a
simple example, the hydrogen molecule. The two nuclei of the hydrogen molecule are strongly held
together, about 0.74 apart, by their covalent bond. This bond consists of the electrostatic attraction
for the nuclei by the shared electrons (balancing the mutual repulsion of the protons). The bond
energy (the energy required to split H2 into H and H) is 432 kJ mole-1. The bond is created by overlap
of the two 1s orbitals. As pictured at right the electrons occupy the space around both nuclei with
their motion largely concentrated in the space between the nuclei.
The electron waves resonate between the two nuclei. Before bonding, there were two 1s atomic
orbitals, and after bonding there will be two molecular orbitals. The bond exists when the electron
pair is located in the one of lower energy of these two molecular orbitals, a symmetric, bonding
orbital.
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The other molecular orbital as shown in the molecular orbital diagram of H2 at right, is of higher
energy, and is antisymmetric and antibonding. This orbital exists along with the bonding orbital, but
the space occupied by the antibonding orbital is beyond the internuclear space (not between the
nuclei). If both electrons to occupy such an orbital, the hydrogen nuclei would both repel one another
and be attracted to the electrons. Imagine a helium atoms, in which the antibonding orbital would be
filled in addition to the bonding orbital, and you can see why two helium atoms show only a weak
attraction.
The Octet Rule: In forming compounds, atoms gain, lose, or share electrons to give a stable electron
configuration characterized by 8 valence electrons.
The octet rule is the most useful in cases involving covalent bonds to C, N, O, and F.
Example: Combine carbon (4 valence electrons) and four fluorines (7 valence electrons each)
The octet rule is satisfied for carbon and each fluorine.
It is common practice to represent a covalent bond by a line. We can rewrite CF4 as
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What’s the driving force for covalent bonding?
• Same thing as in ionic bonding...attaining a complete octet in the outer-most electron shell and
minimizing the energy of the system.
• Atoms are most stable when they have an “octet” of valence electrons (8 electrons in their outer
shell).
– EXCEPTION: Period 1 elements (H, He) have full outer shells when they have 2 valence electrons
(this is called the “duet rule”).
• Group 8 elements (the “noble” gases) don’t form covalent bonds because they already have full
octets….
– EXCEPTION: Xe can use its d-orbitals to form covalent bonds.
If ionic bonding is one limit, and covalent bonding is the other limit, what lies in the middle?
Polar covalent bonds: bonds in which electrons are shared, but the probability distribution is skewed
toward one of the atoms. HF is a typical example.
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The same concept can be envisioned for covalent compounds. Think of the covalent bond as the
electron density existing between the C and H atoms. We can quantify the degree of stabilization by
seeing how much energy it takes to separate a covalent compound into its atomic constituents.
CH4 (g) → C(g) + 4H(g)
ΔH = 1652 kJ/mol
Since we broke 4 C-H bonds with 1652 kJ, the bond energy for a C-H bond is:
1652 kJ/mol
4 bonds = 413 kJ/mol
We can continue this process for a variety of compounds to develop a table of bond strengths (bond
energies).
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We can use these bond energies to determine
Hrxn: H = sum of energy required to break bonds (positive….heat into system) plus the sum of energy
released when the new bonds are formed (negative….heat out from system).
Hrxn = ΣD bonds broken − ΣD bonds formed
(Energy required)
(Energy released)
D is bond energy per mole of bonds
Bond length decreases as bond strength (bond energy) increases.
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LEWIS STRUCTURES
This method of thinking about bonding. The following is a brief review of the rules for Lewis
structures:
1. Normally two electrons pair up to form each bond. This is a consequence of the Pauli exclusion
principle-two electrons must have paired spins if they are both to occupy the same region of space
between the nuclei and thereby attract both nuclei. The definition of a bond as a shared pair of
electrons. however. Is overly restrictive, and we shall see that the early emphasis on electron pairing
in bond formation is unnecessary and even misleading.
2. For most atoms there will be a maximum of eight electrons in the valence shell. This is absolutely
necessary for atoms or the elements lithium through fluorine since they have only four orbitals (an s
and three p orbitals) in I he valence shell. It is quire common, as we do for atoms of other elements to
utilize only their s and p orbitals. Under these conditions the sum of shared pairs (bonds) and
unshared pairs (lone pairs) must equal the number of orbitals-four. This is I he maximum. and for
elements having fewer than four valence electrons, the octet will usually not be filled.
• The order in which the atoms of a molecule are connected is called its constitution or connectivity.
• The constitution of a molecule must be determined in order to write a Lewis structure.
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Lewis electron-dot diagrams, although very much oversimplified, provide a good starting point for
analyzing the bonding in molecules. Credit for their initial use goes to G. N. Lewis,'a n American
chemist who contributed much to thermodynamics and chemical bonding in the early years of the 20 th
century.
In Lewis diagrams, bonds between two atoms exist when they share one or more pairs of electrons.
In addition, some molecules have nonbonding pairs (also called lone pairs) of electrons on atoms.
These electrons contribute to the shape and reactivity of the molecule, but do not directly bond the
atoms together. Most Lewis structures are based on the concept that eight valence electrons
(corresponding to s and p electrons outside the noble gas core) form a particularly stable
arrangement, as in the noble gases with s2p6 configurations. An exception is hydrogen, which is stable
with two valence electrons. Also, some molecules require more than eight electrons around a given
central atom. Simple molecules such as water follow the octet rule, in which eight electrons surround
the oxygen atom. The hydrogen atoms share two electrons each with the oxygen, forming the familiar
picture with two bonds and two lone pairs:
Shared electrons are considered to contribute to the electron requirements of both atoms involved;
thus, the electron pairs shared by H and 0 in the water molecule are counted toward both the
8-electron requirement of oxygen and the 2-electron requirement of hydrogen.
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Only valence electrons are important in bonding. Lewis dot structures show valence electrons
surrounding atom. We visualize the four valence orbitals of an atom as the sides of a box. Electrons
are put into orbitals according to Hund’s rule.
Examples:
Be has 2 valence electrons. Therefore Lewis structure is • Be • , N has 5 valence electrons.
Therefore Lewis structure is
••
•N•
•
, Br has 7 valence electrons. Therefore Lewis structure is
C has 4 valence electrons. Therefore Lewis structure is
•
•C •
•
••
• Br • •,
••
GUIDELINES for Lewis Structures (or Electron Dot Diagrams) of Molecules .
Step 1:
The molecular formula and the connectivity are determined by experiment.
• Example:
Methyl nitrite has the molecular formula CH3NO2. All hydrogens are bonded to carbon, and the order
of atomic connections is CONO.
Step 2:
Count the number of valence electrons. For a neutral molecule this is equal to the number of valence
electrons of the constituent atoms.
• Example (CH3NO2):
Each hydrogen contributes 1 valence electron. Each carbon contributes 4, nitrogen 5, and each
oxygen 6 for a total of 24.
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Step 3:
Connect the atoms by a covalent bond represented by a dash.
• Example: Methyl nitrite has the partial structure:
Step 4:
Subtract the number of electrons in bonds from the total number of valence electrons.
• Example: 24 valence electrons – 12 electrons in bonds. Therefore, 12 more electrons to assign.
Step 5:
Add electrons in pairs so that as many atoms as possible have 8 electrons. Start with the most
electronegative atom.
• Example: The remaining 12 electrons in methyl nitrite are added as 6 pairs.
Step 6:
If an atom lacks an octet, use electron pairs on an adjacent atom to form a double or triple bond.
• Example: Nitrogen has only 6 electrons in the structure shown.
All the atoms have octets in this Lewis structure.
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Inorganic Examples
Organic Examples
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A simple procedure for deciding how to place the electrons is as follows:
1. Determine the total number of valence shell electrons from all the atoms (N) that are available to be
distributed in the structure.
2. Multiply the number of atoms present by eight to determine how many electrons would be required
to give an octet around each atom (S).
3. The difference (S – N) gives the number of electrons that must be shared in the structure.
4. If possible, change the distribution of electrons to give favorable formal charges (discussed later in
this chapter) on the atoms.
For CO, the total number of valence shell electrons is 10, and to give octets around two atoms, it
would require 16 electrons. Therefore, 16 – 10 = 6 electrons must be shared by the two atoms. Six
electrons are equivalent to three pairs or three covalent bonds.
For a molecule such as SO2, we find that the number of valence shell electrons is 18 and three
atoms would require 24 electrons to make three octets. Therefore, 24 – 18 = 6, the number of electrons
that must be shared, which gives a total of three bonds between the sulfur atom and the two oxygen
atoms. However, because we have already concluded that each atom in the molecule must have an
octet of electrons around it, the sulfur atom must also have an unshared pair of electrons in addition
to the three pairs that it is sharing. This can be shown as in the structure
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Lewis Structures (or Electron Dot Diagrams) for POLYATOMIC IONS:
A polyatomic ion consists of covalently bonded atoms with an overall charge.
1. Calculate the total number of electrons (e–s) for all atoms
2. Account for # of e–s associated with charge:
– If ion is positively charged, subtract # of electrons from total
– If ion has +2 charge → subtract 2 electrons from total to get the total # of electrons
– If ion is negatively charged, add # of electrons from total
– If ion has –3 charge → add 3 electrons to get the total # of electrons
3. Divide new total by 2 to get total # of electron pairs
4. Surround central atom (will be indicated) with 4 e– pairs, then distribute outer atoms around central
atom.
5. If any atom (except H) does not have an octet, move nonbonding e–s from central atom to a position
b/w atoms, forming double and triple bonds until all atoms have an octet.
6. Put brackets around all the atoms, and put charge on upper right-hand side
– This indicates the charge belongs to entire entity rather than to a single atom in the ion.
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RESONANCE: In many molecules, the choice of which atoms are connected by multiple bonds is
arbitrary. When several choices exist, all of them should be drawn. For example, as shown in figure ,
three drawings (resonance structures) of CO3 – 2 are needed to show the double bond in each of the
three possible CO positions. In fact, experimental evidence shows that all the CO bonds are identical,
with bond lengths (129 pm) between double-bond and single-bond distances (116 pm and 143 pm
respectively); none of the drawings alone is adequate to describe the molecular structure, which is a
combination of all three, not an equilibrium between them. This is called resonance to signify that
there is more than one possible way in which the valence electrons can be placed in a Lewis
structure. Note that in resonance structures, such as those shown for CO3 – 2 in figure, the electrons
are drawn in different places but the atomic nuclei remain in fixed positions.
The species CO3 – 2, NO3 – and SO3 – 2, are isoelectronic (have the same electronic structure). Their
Lewis diagrams are identical, except for the identity of the central atom. When a molecule has several
resonance structures, its overall electronic energy is lowered, making it more stable. Just as the
energy levels of a particle in a box are lowered by making the box larger, the electronic energy levels
of the bonding electrons are lowered when the electrons can occupy a larger space.
Lewis Diagrams for CO3 – 2
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EXPANDED SHELLS: When it is impossible to draw a structure consistent with the octet rule, it is
necessary to increase the number of electrons around the central atom. An option limited to elements
of the third and higher periods is to use d orbitals for this expansion, although more recent
theoretical work suggests that expansion beyond the s and p orbitals is unnecessary for most main
group molecules. In most cases, two or four added electrons will complete the bonding, but more can
be added if necessary. Ten electrons are required around chlorine in ClF3 and 12 around sulfur in SF6 ,
figure below. The increased number of electrons is described as an expanded shell or an expanded
electron count.
There are examples with even more electrons around the central atom, such as IF7 (14 electrons),
[TaF8] – 3 (16 electrons), and [ XeF8] – 2 (18 electrons). There are rarely more than 18 electrons (2 for s,
6 for p, and 10 for d orbitals) around a single atom in the top half of the periodic table, and crowding of
the outer atoms usually keeps the number below this, even for the much heavier atoms with f orbitals
energetically available.
Structures of ClF3 and SF6
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FORMAL CHARGE: Formal charges can be used to help in the assessment of resonance structures
and molecular topology. The use of formal charges is presented here as a simplified method of
describing structures, just as the Bohr atom is a simple method of describing electronic configuration
in atoms. Both of these methods are incomplete and newer approaches are more accurate, but they
can be useful as long as their limitations are kept in mind.
Formal charges can help in assigning bonding when there are several possibilities. This can
eliminate the least likely forms when we are considering resonance structures and, in some cases,
suggests multiple bonds beyond those required by the octet rule. It is essential, however, to remember
that formal charge is only a tool for assessing Lewis structures, not a measure of any actual charge on
the atoms. Formal charge is the apparent electronic charge of each atom in a molecule, based on the
electron-dot structure. The number of valence electrons available in a free atom of an element minus
the total for that atom in the molecule (determined by counting lone pairs as two electrons and
bonding pairs as one assigned to each atom) is the formal charge on the atom:
Formal charge = number of valence electrons – number of unshared
– number of bonds
in a free atom of the element
electrons on the atom
to the atom
In addition,
Charge on the molecule or ion = sum of all the formal charges
Structures minimizing formal charges, placing negative formal charges on more electronegative
(in the upper right-hand part of the periodic table) elements, and with smaller separation of charges
tend to be favored. Three examples, SCN – , OCN –, and CNO –, will illustrate the use of formal
charges in describing electronic structures
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Examples:
SCN –
In the thiocyanate ion, SCN – , three resonance structures are consistent with the electron -dot
method, as shown below. Structure A has only one negative formal charge on the nitrogen atom, the
most electronegative atom in the ion, and fits the rules well. Structure B has a single negative charge
on the S, which is less electronegative than N. Structure C has charges of 2 – on N and 1+ on S,
consistent with the relative electronegativities of the atoms but with a larger charge and greater
charge separation than the first. Therefore, these structures lead to the prediction that structure A is
most important, structure B is next in importance, and any contribution from C is minor.
The bond lengths are consistent with this conclusion, with bond lengths between those of structures
A and B. Protonation of the ion forms HNCS, consistent with a negative charge on N in SCN – . The
bond lengths in HNCS are those of double bonds, consistent with the structure H – N= C=S.
Table of S – C and ,C – N Bond Lengths (pm)
Resonance structures of thiocyanate, SCN –
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OCN – The isoelectronic cyanate ion, OCN – , has the same possibilities, but the larger electron–
egativity of 0 makes structure B more important than in thiocyanate. The protonated form contains
97% HNCO and 3% HOCN, consistent with structure A and a small contribution from B. The bond
lengths in OCN- and HNCO in Table below are consistent with this picture, but do not agree perfectly.
Table,of O – C and C – N bond lengths (pm)
Resonance Structures of Cyanate, OCN –
CNO – The isomeric fulminate ion, CNO – , can be drawn with thee similar structures, but the resulting
formal charges are unlikely. Because the order of electronegativities is C < N < 0, none of these are
plausible structures and the ion is predicted to be unstable. The only common fulminate salts are of
mercury and silver; both are explosive. Fulminic acid is linear HCNO in the vapor phase, consistent
with structure C, and coordination complexes of CNO with many transition metal ions are known with
MCNO structur.
Resonance Structures of Fulminate, CNO –
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Some molecules have satisfactory electron-dot structures with octets, but have better structures with
expanded shells when formal charges are considered. In each cases in figure below, the observed
structures are consistent with expanded shells on the central atom and with the resonance structure
that uses multiple bonds to minimize formal charges. The multiple bonds may also influence the
shapes of the molecules.
Formal Charge and Expanded Shells
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